Chemistry Past Paper - OCR CHEM 111

CHEMISTRY THE CENTRAL SCIENCE CHEM 111 Module: The Central Science Unit 1 Energy: Electrochemistry and Nuclear Unit 2 Chemistry of Engineering Materials Unit 3 Chemistry of the Environment Electrochemistry Electrochemistry is the branch of chemistry that deals with the study of the re...

CHEMISTRY THE CENTRAL SCIENCE CHEM 111 Module: The Central Science Unit 1 Energy: Electrochemistry and Nuclear Unit 2 Chemistry of Engineering Materials Unit 3 Chemistry of the Environment Electrochemistry Electrochemistry is the branch of chemistry that deals with the study of the relationships between electricity and chemical reactions. It includes the study of both spontaneous and nonspontaneous processes. Electrochemical processes are redox (oxidation-reduction) reactions in which the energy released by a spontaneous reaction is converted to electricity or in which electrical energy is used to cause a nonspontaneous reaction to occur. Oxidation States General Rules 1. In a neutral compound all oxidation numbers must add up to zero. 2. In an ion the all oxidation numbers must add up to the charge on the ion. 3. Free elements have an oxidation number of zero (e.g. Na, Fe, H2, O2, S8). 4. Fluorine = -1 5. Group 1 = +1 Group 2 = +2 6. Hydrogen with Non-Metals = +1 Hydrogen with Metals (or Boron) = -1 7. Oxygen = -2 (except with Fluorine or Peroxides) 8. Group 17(7A) = -1 Group 16 (6A) = -2 Group 15 (5A) = -3 Oxidation States and Oxidation-Reduction Reactions Oxidation-reduction (redox) reactions occur when electrons are transferred from an atom that is oxidized to an atom that is reduced. increase in oxidation state decrease in oxidation state Oxidation States and Oxidation-Reduction Reactions Oxidizing Agent Oxidation number of pure element is 0 Reducing Agent Oxidation-Reduction Reactions It is helpful to view the process with regard to each individual reactant, that is, to represent the fate of each reactant in the form of an equation called a half-reaction: oxidation half-reaction : reduction half-reaction : These equations show that Na atoms lose electrons while Cl atoms (in the CL 2 molecule) gain electrons, the “s” subscripts for the resulting ions signifying they are present in the form of a solid ionic compound. Remembering Oxidation and Reduction It is common to remember the difference between oxidation and reduction using one of two mnemonic devices: 1. "LeO says GeR" 2. "Oil Rig" Loses electrons = Oxidation Oxidation is loss Gains electrons = Reduction Reduction is gain Describing Redox Reactions Solution: Solution: Seatwork: Give the oxidation states and identify the oxidant (oxidizing agent) and reductant (reducing agent), reduced substance and oxidized substance. Half Reactions Fe + Cu2+ → Cu + Fe2+ Oxidation: Fe → Fe2+ + 2e− Reduction: Cu2+ + 2e− → Cu Al + Cu2+ → Al3+ + Cu Oxidation: Al → Al3+ + 3e− Reduction: Cu2+ + 2e− → Cu +3 -2 +2 -2 0 +4 -2 Fe2 O3 + CO → Fe + CO2 Oxidation: CO → CO2 Reduction: Fe2 O3 → Fe Half Reactions +7 -2 +4 -2 +4 -2 +6 -2 MnO− 4 + SO2−3 → MnO2 + SO2−4 Oxidation: SO2− 3 → SO 2− 4 Reduction: MnO−4 → MnO2 +6 -2 +4 -2 +3 +6 -2 Cr2 O2− 7 + SO2 → Cr 3+ + SO2−4 Oxidation: SO2 → SO2− 4 Reduction: Cr2 O2− 7 → Cr 3+ 0 +1+5-2 +2 +5-2 +2-2 +1 -2 Cu + HNO3 → Cu(NO3 )2 + NO + H2 O Oxidation: Cu + HNO3 → Cu(NO3 )2 Reduction: HNO3 → NO Balancing Half Reactions (Acidic Solution) Balance each half-reaction for: - atoms of interest - Oxygen (O) atoms by adding H2O - Hydrogen (H) atoms by adding H+ ions - electrons (charge) by adding electrons 2− this reaction is Aqueous (it takes place in water) MnO− 2− 4 + SO3 → MnO2 + SO4 balance oxygen atoms by adding water Oxidation: SO2− 3 → SO 2− 4 H2 O + SO2− 3 → SO 2− 4 + 2H + + 2e − Reduction: MnO− 4 → MnO2 3e− + 4H+ + MnO− 4 → MnO2 + 2H2 O Balancing Half Reactions (Acidic Solution) Balance each half-reaction for: - atoms of interest - Oxygen (O) atoms by adding H2O - Hydrogen (H) atoms by adding H+ ions - electrons (charge) by adding electrons Fe + Cu2+ → Cu + Fe2+ Oxidation: Fe → Fe2+ + 2e− Reduction: Cu2+ + 2e− → Cu Balancing Half Reactions (Acidic Solution) Balance each half-reaction for: - atoms of interest - Oxygen (O) atoms by adding H2O - Hydrogen (H) atoms by adding H+ ions - electrons (charge) by adding electrons Al + Cu2+ → Al3+ + Cu Oxidation: Al → Al3+ + 3e− Reduction: Cu2+ + 2e− → Cu Balancing Half Reactions (Acidic Solution) Balance each half-reaction for: - atoms of interest - Oxygen (O) atoms by adding H2O - Hydrogen (H) atoms by adding H+ ions - electrons (charge) by adding electrons Fe2 O3 + CO → Fe + CO2 Oxidation: CO → CO2 H2 O + CO → CO2 + 2H+ + 2e− Reduction: Fe2 O3 → Fe 6e− + 6H+ + Fe2 O3 → 2Fe + 3H2 O Balancing Half Reactions (Acidic Solution) Balance each half-reaction for: - atoms of interest - Oxygen (O) atoms by adding H2O - Hydrogen (H) atoms by adding H+ ions - electrons (charge) by adding electrons MnO− 4 → Mn 2+ 5e− + 8H+ + MnO− 4 → Mn 2+ + 4H O 2 C2 H5 OH → C2 H4 O2 H2 O + C2 H5 OH → C2 H4 O2 + 4H+ +4e− Combining Half Reactions Al + Cu2+ → Al3+ + Cu Oxidation: Al → Al3+ + 3e− Reduction: Cu2+ + 2e− → Cu Al → Al3+ + 3e− 2 Cu2+ + 2e− → Cu 3 2Al+3Cu2+ + 6e− → 2Al3+ + 6e− + 3Cu 2Al + 3Cu2+ → 2Al3+ + 3Cu Combining Half Reactions Cu + Ag+ → Cu2+ + Ag Oxidation: Cu → Cu2+ + 2e− Reduction: Ag+ + e− → Ag Cu → Cu2+ + 2e− Ag+ + e− → Ag 2 Cu + 2Ag+ + 2e− → Cu2+ + 2e− + 2Ag Cu + 2Ag+ → Cu2+ + 2Ag Combining Half Reactions +7 -2 -1 0 +2 MnO−4 + I− → I2 + Mn2+ Oxidation: 2I− → I2 +2e− Reduction: 5e− + 8H+ + MnO− 4 → Mn 2+ + 4H O 2 2I− → I2 + 2e− 5 5e− + 8H+ + MnO− 4 → Mn 2+ + 4H O 2 2 10I− + 10e− + 16H+ + 2MnO− 4 → 5I 2 + 10e − + 2Mn 2+ + 8H2 O 10I− + 16H+ + 2MnO− 4 → 5I 2 + 2Mn 2+ + 8H O 2 Combining Half Reactions H2 O + SO2− 3 → SO 2− 4 + 2H + + 2e − 3 3e− + 4H+ + MnO− 4 → MnO2 + 2H2 O 2 3H2 O + 3SO2− − + − 2− + − 3 + 6e + 8H + 2MnO4 → 3SO4 + 6H + 6e + 2MnO2 + 4H2 O 3SO2− 3 + 2H + + 2MnO− → 3SO2− + 2MnO + H O 4 4 2 2 Balancing Half Reactions (Basic Solution) Balance each half-reaction for: - atoms of interest - Oxygen (O) atoms by adding H2O - Hydrogen (H) atoms by adding H+ ions - electrons (charge) by adding electrons 0 +5 -2 +2 +4 -2 Zn + NO− 3 → Zn2+ + NO2 Oxidation: Zn → Zn2+ Zn → Zn2+ + 2e− Reduction: NO− 3 → NO2 e− + 2H+ + NO− 3 → NO2 + H2 O  Combining Half Reactions Zn → Zn2+ + 2e− e− + 2H+ + NO−3 → NO2 + H2 O 2 Zn + 2e− + 4H+ + 2NO−3 → Zn 2+ + 2e− + 2NO + 2H O 2 2 Zn + 4H+ + 2NO− 3 → Zn 2+ + 2NO + 2H O 2 2 - For each H+ , add one OH− to both sides 4OH− + Zn + 4H+ + 2NO− 3 → Zn 2+ + 2NO + 2H O + 4OH− 2 2 - Combine H+ and OH− to make H2 O Zn + 4H2 O + 2NO− 3 → Zn 2+ + 2NO + 2H O + 4OH− 2 2 - Substract H2 O from both sides if possible Zn + 2H2 O + 2NO− 3 → Zn 2+ + 2NO + 4OH− 2 Electrochemistry Electrochemistry is the branch of chemistry that deals with the study of the relationships between electricity and chemical reactions. It includes the study of both spontaneous and nonspontaneous processes. Electrochemical processes are redox (oxidation-reduction) reactions in which the energy released by a spontaneous reaction is converted to electricity or in which electrical energy is used to cause a nonspontaneous reaction to occur.  Two main ways chemical reactions and electricity interact: 1. Certain chemical reaction can create electricity 2. Electricity can make certain chemical reaction happen that wouldn’t happen otherwise When electrons move through a material, they create an electric current.  Galvanic Cells (Voltaic Cells) Neutral atoms make up solid anode cathode metal. site of oxidation site of reduction Metal ions ussually dissolve in water. spontaneous redox reaction Galvanic Cells (Voltaic Cells)  Cell Potentials under Standard Conditions Why do electrons transfer spontaneously from a Zn atom to a Cu2+ ion? Water flows spontaneously over a waterfall because of a difference in potential energy between the top of the falls and the bottom. In a similar fashion, electrons flow spontaneously through an external circuit from the anode of a voltaic cell to the cathode because of a difference in potential energy. The potential energy of electrons is higher in the anode than in the cathode. Thus, electrons flow spontaneously toward the electrode with the more positive electrical potential. Water analogy for electron flow. Just as water spontaneously flows downhill, electrons flow spontaneously from the anode to the cathode in a voltaic cell.  Cell Potentials under Standard Conditions Potential Difference - difference in potential energy per electrical charge between two electrodes - measured in volts (V) one electron has a charge of 1.60 * 10-19 C Cell Potential ( Ecell ) - potential difference between the two electrodes of a voltaic cell - electromotive force (emf) “causing electron motion” Standard cell potential or standard emf (E°cell ) - cell potential under standard condition Standard Reduction Potentials Cell Diagram Zn (s) + Cu2+ (aq) Cu (s) + Zn2+ (aq) Zn (s) | Zn2+ (1 M) || Cu2+ (1 M) | Cu (s) anode cathode STANDARD REDUCTION POTENTIALS AT 25 °C STANDARD REDUCTION POTENTIALS AT 25 °C